Acid mine drainage (AMD) is a common problem for mining industries throughout the world. AMD drainage from metal mining typically contains dissolved metals of high concentration and more than 3 g/L sulfate. The high acidity and presence of these metals make AMD treatment a major concern because of the possible deleterious effects of the effluent on the environment.
There are more than 40,000 remote abandoned mines and a large number of pit lakes in the state of Montana alone. There are many thousands of such mines in other states, including Pennsylvania, Ohio and West Virginia. Acid mine water, upwelling from these remote mines, mainly during the spring season, results in massive destruction of surrounding vegetation. The Berkeley Pit in Butte, Mont., contains 30 billion gallons of acid mind drainage, with daily increases of 3 million gallons per day. This represents a large source of recoverable metals from mine drainage.
Metals are an integral part of the world economy. The residual effects of metals and their use, particularly in aqueous streams, is a continuous problem for metal producers and users, as well as federal and state regulators. Innovating and alternative techniques that allow for the economic control or recovery of metals is one alternative that lends itself not only to human health and environmental protection, but also to resource conservation and reuse of valuable commodities.
Heavy metals can create environmental hazards and are a major pollution (problem for streams that receive acid mine drainage. Metals also appear in wastewaters from metal finishing and metal production facilities, chemical cleaning wastes, as well as ash-pond effluents from coal-fired power plants (Bhattacharyya, 1979).
Treatment of acid mine drainage is a major environmental issue for the mining industries. Old abandoned mines produce acid mine drainage that causes billions of dollars of damage to natural vegetation, lumber trees, rivers, natural habitats and aquatic life. The flow rate of acid water, generated from water introduced by spring thaws and rain, may vary from a few gallons per minute to several thousand gallons per minute. In the U.S., acid mine drainage and other toxins from abandoned mines have polluted 180,000 acres of reservoirs and lakes and 12,000 miles of streams and rivers, (Kleinman, 1989). It has been estimated that cleaning up these polluted waterways would cost U.S. taxpayers between $32 billion and $72 billion. The U.S. Bureau of Mines has estimated that the U.S. mining industry spends over $1 million each day to treat acidic mine water (Pearse, 1996). As noted above, one of the largest locations of acid mine water is the Berkeley Pit in Butte, Mont., encompassing over one square miles in surface area and over 900 feet deep. The water in the Pit has a pH that varies between 2.2 and 2.7. Approximately 3 million gallons of water flow into the Pit daily, resulting in a rise of about 10 feet every nine months.
Many hydrometallurgical processes are based on the solubility behavior of metals in aqueous solutions. Precipitation of metal hydroxides is most easily controlled by pH adjustment and is one of the best known and widely used methods for removing certain metals from impure streams. The conventional approach is to use a base such as lime or sodium hydroxide to raise the pH and precipitate the metals from solution. However, several researchers have used the lower solubilities of metal sulfides to improve metal waste treatment.
Further, sulfide precipitation is becoming more prevalent because lower metal concentrations can be achieved. In many mining operations, however, sulfur compounds may be present, so aqueous metal and sulfide reactions must be considered as well.
The following section mathematically describes the reactions that occur as well as some documented applications of sulfide precipitation reported in the literature.
Monhemius (1977) computed metal concentrations in solution as a function of pH in the presence of hydroxide and sulfide ions using the solubility product of various metal salts.
For hydroxide salts, the pH is important because the hydroxide concentration is limited by the dissociation constant of water, Kw Thus, the concentration of a metal hydroxide in solution can be given as a function of pH.
                              log          ⁢                      {                          M                              Y                +                                      }                          =                                            1              x                        ⁢            log            ⁢                                                  ⁢                          K              Sp                                -                                    Y              x                        ⁢                          (                              pH                +                                  log                  ⁢                                                                          ⁢                                      K                    W                                                              )                                                          (        1        )            
Because hydrogen sulfide will dissociate, its solubility can be calculated using the pH, partial pressure, and dissociation constant. Thus,log {S2−}=2 pH+log PH2S+log KPO  (2)
A parameter pS can be defined as follows:pS=−log {S2−}  (3)
Monhemius uses a dissociation reaction ofH2S→2H++S2− withKPO=[H+]2.[S2−]=1.23×10−23  (4)
Oxtoby and Nachtrieb (1990) present a second, though similar, way to calculate solubility of metal sulfides. They use the fact that the sulfide ion is highly unstable in solution, and propose the following overall reaction for the dissolution of metal sulfides in solution:MS(s)+H2O→M2+(aq)+OH−(aq)+HS−(aq).
The hydroxide ion concentration in solution is fixed by the pH.Kw=[OH−][H+]=10−14  (5)
The concentration of the [HS−] ion is then computed from the acid ionization of H2S.H2S(aq)+H2O→H3O+(aq)+HS−(aq) 
                              K          a                =                                                            [                                                      H                    3                                    ⁢                                      O                    -                                                  ]                            ⁢                                                          [                              HS                -                            ]                                      [                                                H                  2                                ⁢                S                            ]                                =                      9.1            ×                          10                              -                8                                                                        (        6        )            
The solubility of H2S in pure water is given by Morse et al. (1987) as
                                          K            O                    ⁡                      (                                          mol                /                L                            ⁢                              -                            ⁢              atm                        )                          =                              -            41.0563                    +                      66.4005            ⁢                                                  [                          1              7                        ]                    -                      15.1060            ⁢                                                  ⁢                          ln              ⁢                                                          [                              1                7                            ]                                                          (        7        )            
This assumes a fugacity of H2S of one atmosphere. The concentration at other fugacities is also given as:C*(mol/L)=KOfH2S  (8)
Finally, the concentration of the metal is determined by the solubility product.log[M+]=log[KSO]−log[OH−]−log[HS−]  (9)
It is possible to generate a graphical display of the solubility curves from these equations where metal concentrations are determined as a function of pH. Monhemius (1977) has published the solubility curves of four metal sulfides as a function of pH at an H2S pressure of 1 atm at 25° C. using his data and equations. Likewise, similar, but not identical figures can be generated using the data and method of Oxtoby and Nachtrieb (1990). Both methods reveal that there is a specific pH for each metal above which the metal will precipitate out of solution.
Table I gives a list of solubility products for several metal sulfides and hydroxides.
This approach is useful in calculating equilibrium values, but, unfortunately, it has two limitations. One is that it does not account for metal complexes that may form. The second is that the thermodynamic data do not include information about reaction rates. Moreover, while these tables are useful for describing the relationship for a single metal, they are not accurate for a complex ionic solution because they do not account for the “common ion” effects. Therefore, the relationships necessary for process design cannot be predicted theoretically, but rather must be determined experimentally.
TABLE 1Solubility Products at 25° C. (Monhemius, 1977)Log KsoHydroxide(Monhemius,(Sulfide Oxtoby andMetal(Moffliernius, 1977)1977)Sulfide Nachtrieb, 1990)Al3+−32.0——Ca2+−5.3——Cd2+−14.3−28.9−27.2CO2+−14.5−22.1—Cr3+−30.0——CU2+−19.8−35.9−36.3Fe2+−16.3−18.8−18.3Fe3+−38.6——Mg2+−11.3——Mn2+−12.7−13.3−13.5Ni2+−15.3−21.0—Zn2+−16.1−24.5−24.7
Bhattacharyya and co-workers (1981) studied arsenic and heavy metal removal from non-ferrous smelters by controlled precipitation with sodium sulfide and lime in a single stage precipitator. They used a bench-scale process at the University of Kentucky and a full-scale treatment facility (200 m3/hr) at a Swedish copper and lead smelting plant. For the bench-scale process, they used actual scrubber wastewater from a non-ferrous smelting plant. First, the pH of the water was raised to a range of 4.0 to 5.5 using a lime slurry. Then, sodium sulfide was added, polymer was added for sedimentation, and post-filtration removed the sulfide precipitate. The results showed that they could remove 98% of the cadmium, copper, iron (total), selenium, and zinc initially present, with the optimum conditions being a pH of 8.0 and 60% of the theoretical sulfide dosage required. This was possible because some metals, such as copper, have a low solubility at a high pH. The results also indicated that arsenic removal was dependent on the ratio of ferric iron to arsenic (Fe(III)/As ratio) so that, between Fe (III)/As ratios of 0.6 to 2.0, the arsenic removal was greater than 90%. At the Swedish facility, they found that Cd, Cu, Hg and Pb were completely removed by sulfide precipitation. However, the facility was not able to adequately maintain the pH or sulfide dosage (pH values ranged from 3–5; sulfide dosages ranged from 0.8 to 3.1 times the theoretical value needed), and neither zinc nor arsenic was removed. Their research showed that arsenic sulfide precipitates better at a pH below 3, whereas zinc precipitates better at a pH above 5. Also, they found that dissolved sulfite (SO22−) present in the water consumed some of the sulfide, thus reducing the amount of sulfide available for precipitation.
Previous processes, as described in many patents and publications, have attempted to remove metals from acidic waste streams to minimize the environmental impact of the wastewater release. Such processes are designed to remove all contaminants in a single stage, or as few stages as possible, with the result that the precipitated metals are co-mingled. These precipitates have little or no commercial value and are usually treated as a waste material.